Ethylene Glycol

   Ethylene glycol (ETH-uh-leen GLYE-kol) is clear, colorless,
syrupy liquid with a sweet taste. One should not attempt
to confirm the compound’s taste, however, as it is toxic. In
recent years, more than 4 billion kilograms (9 billion pounds)
of ethylene glycol has been produced in the United States
annually. The compound is used primarily as an antifreeze
and in the manufacture of a number of important chemical
compounds, including polyester fibers, films, bottles, resins,
and other materials.
   Ethylene glycol was first prepared in 1859 by the
French chemist Charles Adolphe Wurtz (1817–1884).
Wurtz’s discovery did not find an application, however,
until the early twentieth century, when the compound was
manufactured for use in World War I (1914–1918) in the
manufacture of explosives and as a coolant. By the 1930s,
a number of uses for the compound had been found, and the
chemical industry began producing ethylene glycol in large
quantities.
   The primary method of producing ethylene glycol
involves the hydration of ethylene oxide, a ring compound
consisting of two methylene (-CH2) groups and one oxygen
atom. Hydration is the process by which water is added to a
compound. The hydration of ethylene oxide is conducted at a
temperature of about 383F (195C) without a catalyst, or at
about 50C to 70C (122F to 158F) with a catalyst, usually a
strong acid, either process resulting in a yield of at least 90
percent of ethylene glycol.
   Other methods of preparation are also available. For
example, the compound can be produced directly from synthesis
gas, a mixture of carbon monoxide and hydrogen; or by
treating ethylene (CH2=CH2) with oxygen in an acetic acid
solution using a catalyst of tellurium oxide or bromide ion.
One of the first major uses of ethylene glycol was as a
radiator coolant in airplanes. The compound actually made
possible a change in the design of airplanes. At one time,
plain water was used as the coolant in airplane radiators. The
faster the airplane flew, the greater the risk that its radiator
would boil over. Adding ethylene glycol to the water raised
the boiling point of the coolant and allowed airplanes to fly
faster with smaller radiators. This change was especially
useful in the construction of military airplanes used in
combat.

   Ethylene glycol is still used extensively as a coolant and
antifreeze in cooling systems. It is also used as a deicing
fluid for airport runways, cars, and boats. Brake fluids and
shock-absorber fluids often contain ethylene glycol as protection
against freezing. About 26 percent of all the ethylene
glycol made in the United States is used for some kind of
cooling or antifreeze application.

   The largest single use of ethylene glycol today is in the
manufacture of a plastic called polyethylene terephthalate
(PET). PET’s primary application is in the manufacture of
plastic bottles, an application that accounts for about a third
of all the ethylene glycol made in the United States. Large
amounts of PET are also used in the manufacture of polyester
fibers and films. Some additional uses of the compound
include:

• As a humectant (a substance that attracts moisture) in
keeping some food, tobacco, and industrial products
dry;
• As a solvent in some paints and plastics;
• In the dyeing of leathers and textiles;
• In the manufacture of printing inks, wood stains, ink
for ball-point pens, and adhesives;
• In the production of artificial smoke and fog for theatrical
productions;
• As a stabilizer in the soybean-based foam sometimes
used to extinguish industrial fires; and
• In the manufacture of specialized types of explosives.
   Ethylene glycol poses a number of potential health and
safety hazards. It is very flammable and highly toxic. Ingestion
of the compound may cause nausea, vomiting, abdominal
pain, weakness, convulsions, and cardiac problems.
Higher doses can result in severe kidney damage that leads
to death.

Amino Acid

amino acid An organic molecule possessing both
acidic carboxylic acid (–COOH) and basic amino
(–NH2) groups attached to the same tetrahedral carbon
atom.
Amino acids are the principal building blocks of
proteins and enzymes. They are incorporated into
proteins by transfer RNA according to the genetic
code while messenger RNA is being decoded by ribo-
somes. The amino acid content dictates the spatial
and biochemical properties of the protein or enzyme
during and after the final assembly of a protein.
Amino acids have an average molecular weight of
about 135 daltons. While more than 50 have been dis-
covered, 20 are essential for making proteins, long
chains of bonded amino acids.
Some naturally occurring amino acids are alanine,
arginine, asparagine, aspartic acid, cysteine, glutamine,
glutamic acid, glycine, histidine, isoleucine, leucine,
lysine, methionine, phenylalanine, proline, serine, thre-
onine, tryptophan, tyrosine, and valine.
The two classes of amino acids that exist are
based on whether the R-group is hydrophobic or
hydrophilic. Hydrophobic or nonpolar amino acids
tend to repel the aqueous environment and are located
mostly in the interior of proteins. They do not ionize
or participate in the formation of hydrogen bonds. On
the other hand, the hydrophilic or polar amino acids
tend to interact with the aqueous environment, are
usually involved in the formation of hydrogen bonds,
and are usually found on the exterior surfaces of pro-
teins or in their reactive centers. It is for this reason
that certain amino acid R-groups allow enzyme reac-
tions to occur.
The hydrophilic amino acids can be further subdi-
vided into polar with no charge, polar with negatively
charged side chains (acidic), and polar with positively
charged side chains (basic).

Chlorophyll : The Green

  Chlorophyll (KLOR-uh-fill) is the pigment that gives
plants, algae, and cyanobacteria their green color. The name
comes from a combination of two Greek words, chloros,
meaning ‘‘green’’ and phyllon, meaning ‘‘leaf.’’ Chlorophyll is
the substance that enables plants to create their own food
through photosynthesis.
At least five forms of chlorophyll exist. They are:
• chlorophyll a (also known as a-chlorophyll), with a formula
of C55H72O5N4Mg
• chlorophyll b (also known as b-chlorophyll), with a formula
of C55H70O6N4Mg
• Chlorophyll c1, with a formula of C35H30O5N4Mg
• Chlorophyll c2, with a formula of C35H28O5N4Mg
• Chlorophyll d, with a formula of C54H70O6N4Mg
Chlorophyll a occurs in all types of plants and in algae.
Chlorophyll b is found primarily in land plants. Chlorophyll
c1 and chlorophyll c2 are present in various types of algae.
Chlorophyll d is found in red algae.

   All forms of chlorophyll have a similar chemical structure.
They have a complex system of rings made of carbon
and nitrogen known as a chlorin ring. The five forms of
chlorophyll differ in the chemical groups attached to the
chlorin ring. These differences result in slightly different
colors of the five chlorophylls.
   French chemists Pierre-Joseph Pelletier (1788–1842) and
Joseph-Bienaime´ Caventou (1795–1877) first isolated chlorophyll
in 1817. In 1865, German botanist Julius von Sachs
(1832–1897) demonstrated that chlorophyll is responsible
for photosynthetic reactions that take place within the cells
of leaves. In the early 1900s, Russian chemist Mikhail Tsvett
(1872–1920) developed a technique known as chromatography
to separate different forms of chlorophyll from each
other. In 1929, the German chemist Hans Fischer (1881–
1945) determined the complete molecular structure, making
possible the first synthesis of the molecule in 1960 by the
American chemist Robert Burns Woodward (1917–1979).

   Plants make chlorophyll in their leaves using materials
they have absorbed through their roots and leaves. The
synthesis of chlorophyll requires several steps involving
complex organic compounds. First, the plant converts a common
amino acid, glutamic acid (COOH(CH2)2CH(NH2)COOH)
into an alternative form known as 5-aminolevulinic acid
(ALA). Two molecules of ALA are then joined to form a ring
compound called porphobilinogen. Next, four molecules of
porphobilinogen are joined to form an even larger ring structure
with side chains. Oxidation of the larger ring structure
introduces double bonds in the molecule, giving it the ability
to absorb line energy. Finally, a magnesium atom is introduced
into the center of the ring and side chains are added to
the ring to give it its final chlorophyll configuration.
   Plants store chlorophyll in their chloroplasts, organelles
(small structures) that carry out the steps involved in photosynthesis.
Each chloroplast contains many clusters of several
hundred chlorophyll molecules called photosynthetic units.
When a photosynthetic unit absorbs light energy, chlorophyll
molecules move to a higher energy state, initiating
the process of photosynthesis. The overall equation for the
process of photosynthesis is 6CO2 + 6H2O ! C6H12O6 + 6O2.
That simple equation does not begin to suggest the complex
nature of what happens during photosynthesis. Botanists
divide that process into two major series of reactions: the light
reactions and the dark reactions. In the light reactions, plants
use the energy obtained from sunlight to make two compounds,
adenosine triphosphate (ATP) and nicotinamide adenine
dinucleotide phosphate (NADPH). ATP and NADPH are
not themselves components of carbohydrates, the final products
of photosynthesis. Instead, they store energy that is
used to make possible a series of thirteen different chemical
reactions that occur during the dark stage of photosynthesis
that result in the conversion of carbon dioxide and water
to the simple carbohydrate glucose (C6H12O6).

Riboflavin: The Vitamin B2

 

 

Riboflavin (REY-bo-FLAY-vin), commonly known as vitamin
B2, is an orange-yellow crystalline solvent with a bitter
taste. It is relatively stable when exposed to heat, but tends
to decompose in the presence of light for extended periods of
time. Riboflavin is used in the body for a variety of functions,
including the metabolism of carbohydrates for the
production of energy and the production of red blood cells.
Riboflavin was found in 1879 by Alexander Wynter Blyth (1844-1921) who noticed a compound in cow’s milk that glowed with a yellow fluorescence
when exposed to light. Blyth called the compound
lachtochrome (lachto- = ‘‘milk’’ and -chrome = color), but was
unable to determine its chemical composition or its chemical
properties. In fact, it was not until the 1930s that the
chemical nature of the compound was determined. The
Swiss chemist Paul Karrer (1889–1971) and the Austrian-
German chemist Richard Kuhn (1900–1967) independently
determined the chemical structure of riboflavin and first
synthesized the compound. The name riboflavin is derived
from the fact that the vitamin was first found in association
with the sugar ribose.

   Naturally, plants and microorganisms can synthesize riboflavin.
Some foods rich in riboflavin are brewer’s yeast, dark
green vegetables, mushrooms, legumes, nuts, milk and other
dairy products, sweet potatoes, and pumpkins. Bacteria that
live in the human digestive tract are also able to synthesize
some riboflavin, but not enough to meet the body’s requirement
for the vitamin.
   Riboflavin is produced synthetically using either the
genetically-modified bacterium Bacillus subtilis or a fungus
called Ashbya gossifyii. The bacteria or fungus are cultured
in a large vat that has been seeded with small amounts of
riboflavin. Over time, the organisms generate large quantities
of riboflavin until some desired amount of the compound
has been produced. The vat is then heated to a
temperature sufficient to kill the bacteria or fungi, leaving
crystalline riboflavin behind. The riboflavin is then separated
and purified.

   The human body needs riboflavin to use oxygen efficiently
in the metabolism of amino acids, fatty acids, and
carbohydrates. The vitamin is involved in the synthesis of
niacin (another B vitamin), it activates vitamin B6, and it
helps the adrenal gland to produce hormones. It helps the
body make antibodies to fight disease and infection, regulates
the thyroid gland, and is important in maintaining healthy
hair, nails, and skin. Riboflavin is especially important during
periods of rapid growth because it is involved in the formation
and growth of cells, especially red blood cells.
The human body needs riboflavin to use oxygen efficiently
in the metabolism of amino acids, fatty acids, and
carbohydrates. The vitamin is involved in the synthesis of
niacin (another B vitamin), it activates vitamin B6, and it
helps the adrenal gland to produce hormones. It helps the
body make antibodies to fight disease and infection, regulates
the thyroid gland, and is important in maintaining healthy
hair, nails, and skin. Riboflavin is especially important during
periods of rapid growth because it is involved in the formation
and growth of cells, especially red blood cells.
most likely to suffer from riboflavin deficiency problems are
those with anorexia (a condition in which people refuse to
eat adequate amounts of food), older people with poor diets,
alcoholics (because alcohol impairs a person’s ability to
absorb and use the vitamin), and newborn babies being treated
for jaundice by exposure to ultraviolet light (because
light destroys riboflavin).

Dimethyl Ketone

Dimethyl ketone (DYE-meth-el KEY-tone) is a clear, colorless,
highly volatile and highly flammable liquid with a
characteristic sweet odor and taste. The compound is almost
universally known in chemistry laboratories and industrial
applications by its common name of acetone.
  Acetone was apparently first prepared in 1610 by the
French alchemist Jean Be´guin (c. 1550–c. 1650). Be´guin
obtained acetone by heating lead acetate (also known as
Saturn’s salt) to a high temperature. He obtained a sweetsmelling,
very flammable liquid that he named ‘‘burning
spirit of Saturn.’’ One of the first uses to which the substance
was put was as a solvent in the extraction of the active
constituents of opium. In 1833, the French chemist Antoine
Bussy (1794–1882) gave the compound its modern name of
acetone. The correct chemical formula for acetone was determined
independently in 1832 by the French chemist Jean
Baptiste Andre´ Dumas (1800–1884) and the German chemist
Justus von Liebig (1803–1873).
 
Most of the acetone produced today is made by one of
four methods:
• In the Hock process, cumene [C6H5CH(CH3)2] is first
oxidized to produce cumene hydroperoxide
[C6H5C(CH3)2COOH], which is then reduced to produce
acetone and phenol (C6H5OH); or
• Isopropyl alcohol (2-propanol; CH3CHOHCH3) is oxidized
over a catalyst to obtained acetone; or
• Butane (C4H10) is oxidized to obtain acetone; or
• Acetone is obtained as a by-product of the manufacture
of glycerol [C3H5 (OH)3].

  Acetone’s primary applications are based on its ability to
dissolve such a wide array of organic substances. It is used as
a solvent for paints, varnishes, lacquers, inks, glues, rubber
cements, fats, oils, waxes, and various types of rubber and
plastics. It is perhaps best known to the average person as
the primary ingredient in nail polish remover. The largest
single use of the compound is as a raw material in the
manufacture of other organic chemicals, such as chloroform,
acetic acid, iodoform, bromoform, isoprene, rayon, and photographic
film. It also finds application in storing acetylene
gas (because it absorbs up to 24 times its own weight of the
gas), to clean and dry chemical equipment and electronic
parts, and for the extraction of components of plant and
animal tissues.
  The primary safety concern about acetone is its extreme
flammability. Workers who handle the compound must use
great care to prevent its coming into contact or even being in
the vicinity of open flames. Under the proper conditions,
acetone is also explosive. Exposure of the skin, eyes, and
respiratory system to acetone may produce mild symptoms,
such as dizziness, headaches, and disorientation and irritation
of the eyes and skin. Such conditions are rare, however,
and no long-term health effects of the compound have as yet
been discovered.

Retinol: The Vitamin A

  Retinol (RET-uh-nol) is the scientific name for vitamin A,
a vitamin found only in animals. It occurs as a yellowish to
orange powder with a slight brownish cast and is a relatively
stable compound. Retinol is converted in the body from an
alcohol to the corresponding aldehyde, retinal (C20H28O), one
of the primary chemical compounds involved in the process
by which light is converted to nerve impulses in the retina of
the eye. Vitamin A is also required for a number of other
biochemical reactions in the body, including growth and
development of tissue and maintenance of the immune system

  Vitamin A is synthesized in animal bodies through a
variety of pathways. One important source of vitamin A is a
group of related compounds called the carotenes, substances
responsible for the yellowish or orangish appearance of
fruits and vegetables such as carrots, sweet potatoes, squash,
cantaloupe, apricots, pumpkin, and mangos. Some leafy
green vegetables, such as collard greens, spinach, and kale,
are also good sources of the carotenes. The most important of
the carotenes is b-carotene (beta-carotene), C40H56. The oxidation
of carotenes in animal bodies converts them to retinol.

  The chemical structure of retinol was determined in 1931
by Swiss chemist Paul Karrer (1889–1971), and the compound
was first prepared synthetically shortly thereafter by Austrian-
German chemist Richard Kuhn (1900–1967). The first
successful process for producing retinol commercially was
developed in the mid-1940s by German chemist Otto Isler
(1920–1992), then employed at the pharmaceutical company
Roche, located in Sissein, Germany. Isler’s process involved a
complex series of reactions that begins with the combination
of a fourteen carbon hydrocarbon and a six carbon hydrocarbon
to create the fundamental backbone from which the
retinol molecule is constructed. Regular production of vitamin
A began in 1948 with a projected output of 10 kilograms
per month, which before long was raised to 50 kilograms per
month. The Roche plant at Sissein continues to produce
retinol today.
  Vitamin A is probably best known for its role in maintaining
normal vision. Deficiencies of the compound are
likely to manifest themselves earliest in a variety of eye
problems, most commonly night blindness. Night blindness
is a condition in which one loses the ability to distinguish
objects in reduced light. If left untreated, vitamin A deficiencies
may lead to decreased ability to see in normal light and,
eventually, to complete blindness.
  But vitamin A has been shown to have a number of other
functions in the body. It is essential for the maintenance of
growth, bone formation, reproduction, proper immune system
function, and healing of wounds. A number of additional
claims have been made for the compound, although evidence
is not as strong as it is for the above functions. For example,
it may be effective in preventing or treating a variety of
conditions such as measles, intestinal parasites, osteoporosis,
inflammatory bowel disease, bone marrow disorders, certain
types of cancer, tuberculosis, peritonitis, osteoarthritis, food
poisoning, Alzheimer’s disease, miscarriage, and HIV/AIDS.
In each of these cases, evidence is not yet strong enough to
show a clear-cut connection between retinol and disease, but
research is being conducted to determine how strong the
association may be.
  Retinol is available commercially in a variety of formulations,
including tablets, capsules, and creams. Such products
usually contain a modified form of retinol that is more easily
absorbed by the body. For example, a product known as
tretinoin is a synthetic form of retinol known as all-trans
retinoic acid. The term all trans means that all of the double
bonds in retinoic acid are located on the same side of the
molecule. Products containing tretinoin are used to treat
acne, pimples, wrinkles, blackheads, freckles, sun-spots, and
even pre-cancerous lesions. They work by increasing the rate
with which the skin sheds old cells and replaces them with
new cells.
  Vitamin A supplements in pill or capsule form are available
in two formulations, those that contain retinol and
those that contain beta carotene. It is not possible to take
too much of the latter type of vitamin A. The body will not
convert excess amounts of carotene into retinol but will,
instead, excrete the excess in the urine or stool. An excess
of retinol-based vitamin A, by contrast, may result in certain
medical problems. Since the vitamin is fat soluble, in
may be stored in body fat and reach relatively high concentrations
if too much is ingested. An excess of retinol in the
body may be associated with liver damage, osteoporosis,
rash, fatigue, bone and joint pain, nausea, insomnia, and
personality changes.

Hydrogen Chloride

  Hydrogen chloride (HY-druh-jin KLOR-ide) is a colorless
gas with a strong, suffocating odor. The gas is not flammable,
but is corrosive, that is, capable of attacking and reacting
with a large variety of other compounds and elements.
Hydrogen chloride is most commonly available as an aqueous
solution known as hydrochloric acid. It is one of the most
important industrial chemicals in the world. In 2004, just
over 5 million metric tons (5.5 million short tons) of hydrogen
chloride were produced in the United States, making it
the eighteenth most important chemical in the nation for
that year.
  Hydrogen chloride has probably been known as far back
as the eighth century, when the Arabian chemist Jabir ibn
Hayyan (c. 721–c. 815; also known by his Latinized name of
Geber) described the production of a gas from common table
salt (sodium chloride; NaCl) and sulfuric acid (H2SO4). The
compound was mentioned in the writings of a number of
alchemists during the Middle Ages and was probably first

produced in a reasonably pure form by the German chemist
Johann Rudolf Glauber (1604–1670) in about 1625. The first
modern chemist to prepare hydrogen chloride and describe
its properties was the English chemist Joseph Priestley
(1733–1804) in 1772. Forty years later, in 1818, the English
chemistry and physicist Humphry Davy (1778–1829) showed
that the compound consisted of hydrogen and chlorine, giving
it the correct formula of HCl.
  Commercial production of hydrogen chloride had its
beginning in Great Britain in 1823. The method of production
most popular there and, later, throughout Europe was
one originally developed by the French chemist Nicholas
Leblanc (1742–1806) in 1783. Leblanc had invented the process
as a method for producing sodium hydroxide and sodium
carbonate, two very important industrial chemicals. Hydrogen
chloride was produced as a byproduct of the Leblanc
process, a byproduct for which there was at first no use.
The gas was simply allowed to escape into the air. The suffocating
and hazardous release of hydrogen chloride prompted
governments to pass legislation requiring some other means
of disposal for the gas. In England, that law was called the
Alkali Act and was adopted by the parliament in 1863. Unable
to release hydrogen chloride into the air, manufacturers
began dissolving it in water and producing hydrochloric acid.
Before long, a number of important commercial and industrial
uses for the acid itself were discovered. The ‘‘useless’’
byproduct of the Leblanc process soon became as important
as the primary products of the process, sodium hydroxide
and sodium carbonate.

  Hydrogen chloride is still sometimes made today by the
traditional process of reacting sodium chloride (NaCl) with a
sulfate, such as sulfuric acid or iron(II) sulfate (FeSO4). However,
more than 90 percent of the hydrogen chloride produced
throughout the world today comes as the byproduct of
the chlorination of organic compounds. Chlorination is the
process by which chlorine gas reacts with an organic compound,
usually replacing some of the hydrogen present in the
compound. Since a large number of important chlorinated
organic compounds are produced each year, large amounts of
hydrogen chloride gas are produced as a byproduct. That gas
is simply removed from the reaction and stored in cylinders
for future use. Other methods of producing hydrogen chloride
include the direct synthesis of hydrogen gas and chlorine
gas (producing a very pure product) and the reaction of
sodium chloride, sulfur dioxide, oxygen, and water with each
other at high temperatures (the Hargreaves process).

  Hydrogen chloride and hydrochloric acid have some uses
in common, and some that are different from each other. In
both dry and liquid form, the largest single use of hydrogen
chloride is in the synthesis of organic and inorganic chlorides.
A large number of compounds important in commerce
and industry contain chlorine, including most pesticides,
many pharmaceuticals, and a number of polymeric products.

  Hydrochloric acid is also used widely in the processing of
metallic ores and the pickling of metals. Pickling is the
process by which a metal is cleaned, usually with an acid,
to remove rust and other impurities that have collected on
the metal. Some additional uses of hydrogen chloride and
hydrochloric acid include the following:
• In the brining of foods and other materials. Brining is
the process by which a material is soaked in a salt
solution, usually in order to preserve the material;
• In the treatment of swimming pool water;
• As a catalyst in industrial chemical reactions;
• In the manufacture of semiconductors and other electronic
components;
• To maintain the proper acidity in oil wells (to keep oil
flowing smoothly);
• For the etching of concrete surfaces;
• In the production of aluminum, titanium, and a number
of other important metals.
  Both hydrogen chloride and hydrochloric acid pose serious
health risks to humans and other animals. The gas is an
irritant to the eyes and respiratory system, causing coughing,
choking, and tearing, as well as more serious damage to tissues.
Hydrochloric acid can burn the skin and mucous membranes.
Exposure of only five parts per million of the gas can
produce noticeable symptoms of distress, and exposure of
more than 2,000 parts per million can be fatal. If hydrochloric

acid gets into the eyes, blindness may result. Since hydrochloric
acid is present in many household products, users should
exercise great care when working with such materials.

Nitrous Oxide

 

  Nitrous oxide (NYE-truss OX-side) is also known as
dinitrogen oxide, dinitrogen monoxide, nitrogen monoxide,
and laughing gas. It is a colorless, nonflammable gas with a
sweet odor. Its common name of laughing gas is derived from
the fact that it produces a sense of light-headedness when
inhaled. The gas is widely used as an anesthetic, a substance
that reduces sensitivity to pain and discomfort.
  Nitrous oxide was probably first produced by the English
chemist and physicist Robert Boyle (1627–1691), although he
did not recognize the new compound he had found. Credit for
the discovery of nitrous oxide is, therefore, usually given to
the English chemist Joseph Priestley (1733–1804), who produced
the gas in 1772 and named it ‘‘nitrous air.’’ Other early
names used for the gas include ‘‘gaseous of azote’’ (nitrogen)
and ‘‘oxide of speton.’’ The most complete experiments on the
gas were conducted by the English chemist and physicist Sir
Humphry Davy (1778–1829), who tested nitrous oxide on
himself and his friends. He found that the gas could lessen

pain and discomfort and provided a sense of relaxation and
well-being. Before long, doctors were making use of Davy’s
discovery by using nitrous oxide as an anesthetic.
  The public found other uses for the gas as well. During
the Victorian period in England, members of the upper class
often held laughing gas parties at which people gathered to
inhale nitrous oxide as a recreational drug, rather than for
any therapeutic purpose. In the United States, the showman
P. T. Barnum (1810–1891) created a sideshow exhibit in which
people were invited to test the effects of inhaling nitrous
oxide. After seeing a demonstration of this kind, the American
dentist Horace Wells (1815–1848) first used nitrous oxide
as an anesthetic on his patients.
  In 1868, the American surgeon Edmund Andrews (1824–
1904) extended the use of nitrous oxide as an anesthetic for
his surgical patients. He mixed the gas with oxygen to
ensure that patients received enough oxygen while receiving
the anesthetic. The gas is still widely used by dentists as a
safe and relatively pleasant way of helping patients endure
the discomfort of drilling and other dental procedures.

  The most common commercial method of producing
nitrous oxide involves the controlled heating of ammonium
nitrate (NH4NO3). The compound decomposes to form nitrous
oxide and water. The reaction is essentially the same one
originally used by Priestley in 1772. Although an efficient
means of producing the gas, the reaction must be carried out
with extreme care as ammonium nitrate has a tendency to
decompose explosively when heated. Nitrous oxide can also
be produced by the decomposition of nitrates (compounds
containing the NO3 radical), nitrites (compounds containing
the NO2) radical, or nitriles (compounds containing the CH
radical).

  Nitrous oxide is best known and most widely used as an
anesthetic. Its use is limited primarily to dental procedures
and minor surgeries. Dentists favor nitrous oxide as an anesthetic
because the gas does not make patients completely
unconscious and does not require an anesthesiologist to
administer it. Nitrous oxide works as an anesthetic by blocking
neurotransmitter receptors in the brain, preventing pain
messages from being transmitted.
  Nitrous oxide is also used as a fuel additive in racing
cars, in which case it is often referred to as nitro. The gas is
injected into the intake manifold where it mixes with air and
fuel vapors. Since it breaks down at the high temperatures in
the car’s engine, it provides additional oxygen to increase the

efficiency with which the fuel burns. During World War II,
pilots used nitrous oxide for a similar purpose in their airplanes.
Some additional uses of nitrous oxide include:
• As a propellant in food aerosols;
• For the detection of leaks;
• As a packaging gas for potato chips and other snack
foods, preventing moisture from making the product
become stale;
• In the preparation of other nitrogen compounds; and
• As an oxidizing agent for various industrial processes.
  Nitrous oxide is safe to use in moderate amounts under
controlled conditions. Some people use the compound as a
recreational drug, however, hoping to get a ‘‘high’’ from inhaling
it. One risk of this practice is that the inhalation of
nitrous oxide may reduce the amount of oxygen a person
receives. Also, some long-term health effects, such as anemia
(low red blood cell count) and neuropathy (damage to the
nerves), have been associated with excessive use of the compound.
The use of nitrous oxide for recreational purposes is a
crime in some states.

Carbon Dioxide

  Carbon dioxide (KAR-bun dye-OK-side) is a colorless,
odorless, tasteless, non-combustible gas that can also exist
under pressure as a clear, colorless, odorless, tasteless liquid
and as a white, snow-like solid commonly known as dry ice.
When dry ice is warmed it sublimes (passes directly from
the solid to the gaseous state without first melting) at
78.4C (-109F).
  The true nature of carbon dioxide was discovered over an
extended period of time beginning with the research of the
Flemish physician and chemist Jan Baptista van Helmont
(1580–1635?). In about 1603, van Helmont isolated a gas
produced during the combustion of wood and proved that it
was distinct from air. At the time, air was generally regarded
as an element that could not be divided into separate components.
Van Helmont called the gas gas sylvestre (‘‘wood gas’’),
a substance we now know to be carbon dioxide. Credit for
understanding the true nature of carbon dioxide also goes to
the Scottish chemist Joseph Black (1728–1799) who produced

carbon dioxide by heating calcium carbonate (CaCO3). Black
called the gas fixed air and conducted the first extensive
studies of its properties.
  The first practical use for carbon dioxide was discovered
in the mid-eighteenth century by the English chemist Joseph
Priestley (1733–1804). Priestley found that passing carbon
dioxide into water produced a sparkling, refreshing drink
that he predicted would one day become a great commercial
success. He was, of course, correct, since water containing
carbon dioxide is the basic component of which all soda
drinks are made.

  Carbon dioxide is produced in nature by a number of
reactions. Among the most common is the combustion
(burning) of the fossil fuels (coal, oil, and natural gas). The
gas is also produced during the decay of organic material,
the fermentation of carbohydrates by yeast, and the respiration
of animals. In the laboratory, the simplest and most
direct method of preparation is to treat a carbonate, such
as calcium carbonate, with an acid, such as hydrochloric
acid (HCl).
  Carbon dioxide is obtained commercially as the byproduct
of a number of industrial reactions. For example,
when calcium carbonate is heated to produce lime (CaO),
carbon dioxide is released and captured as a by-product. The
steam reforming (refining) of petroleum results in the production
of a mixture of gases known as synthesis gas, consisting
of carbon dioxide, carbon monoxide, hydrogen, and
nitrogen. Carbon dioxide can be separated from the other
components of synthesis gas for commercial uses. Carbon
dioxide also produces as a by-product of the manufacture of
ammonia (NH3) by the Haber-Bosch process.

  Carbon dioxide plays an essential role in most biological
processes that take place on Earth’s surface. Plants use carbon
dioxide as a raw material to make the carbohydrates on
which their structures are based. When animals eat plants,
those carbohydrates are then used to build and maintain
their body structures.
In addition to its role in natural processes, carbon dioxide
has many commercial and industrial applications. One of the
most important uses is in the carbonation of beverages.
Although beers and sparkling wines contain carbon dioxide
from natural sources (the fermentation of sugars by
yeasts), nearly all carbonated beverages have their carbon
dioxide added artificially. The carbon dioxide adds a zesty
taste to the beverage and helps to preserve it.

  Carbon dioxide is also used as a fire extinguishing agent.
Its use for this purpose is based on the facts that it does not
burn itself and is heavier than air. Thus, when sprayed on a
fire, carbon dioxide settles down on top of the flames and
prevents oxygen from reaching the burning material. The
carbon dioxide can be supplied in a variety of ways in a fire
extinguisher. In some devices, carbon dioxide gas is produced
as the result of a chemical reaction that occurs within
the fire extinguisher. In other devices, liquid carbon dioxide
is released from the extinguisher.
  Carbon dioxide is also used in gaseous, liquid, or solid
form as a refrigerant. As a gas, it is used as the ‘‘working
fluid’’ in refrigerators, the fluid that circulates through the
refrigerator changing back and forth from gas to liquid,
absorbing heat in the process. In the form of dry ice, carbon
dioxide is a very efficient and convenient method for cooling
objects to very low temperatures (close to the sublimation
point of carbon dioxide, about 78.4C (109F).
Some other uses of carbon dioxide include the following:
• As an aerosol propellant;
• To provide an oxygen-free atmosphere in which to conduct
welding and other operations with flammable
materials;
• In the industrial manufacture of carbonates;
• For cloud seeding to promote modifications in the
weather (increases or decreases in rain fall);
• In the fumigation of rice to preserve the product for
extended periods of time;
• As an artificial smoke in theater productions;

• As a moderator to slow down the speed of neutrons
traveling in a nuclear power plant;
• In the frozen food industry;
• To enrich the air in a greenhouse, providing additional
carbon dioxide to promote plant growth; and
• For the hardening of foundry molds and cores.
In general, carbon dioxide poses little or not threat to
humans in concentrations to which one is normally exposed.
Dry ice may pose a hazard if not handled carefully as its very
low temperature can cause damage to the skin.

Ethyl Alcohol

 

  Ethyl alcohol (ETH-uhl AL-ko-hol) is a clear, colorless,
flammable liquid with a sharp, burning taste and a pleasant,
wine-like odor. It is one of the first chemical substances
discovered and used by humans. Ceramic jugs apparently
designed to hold beer have been dated to the Neolithic Period,
about 10,000 BCE. Some scholars suggest that humans
may have learned how to make beer and incorporated it into
their daily diets even before they made and used bread. The
making and use of wine is a clear theme in Egyptian pictographs
dating to the fourth millennium BCE. There probably
does not exist a human culture today in which alcohol consumption
does not occur. Today, beverages with alcohol content
ranging as low as two to five percent (‘‘near beer’’ and
beer) to as high as 50 percent (some forms of vodka) are
known and consumed by humans. In spite of its widespread
use as a beverage, ethyl alcohol has a number of commercial
and industrial uses that account for more than 90 percent of
all the compound produced in the United States.

Ethyl alcohol is made in one of two ways: naturally,
through the process of fermentation, or synthetically, beginning
with compounds found in petroleum. Until the beginning
of World War II, more than 90 percent of all ethyl
alcohol produced in the United States and other developed
nations was made by fermentation. Waste syrup left over
from the production of sugar from sugar cane was treated
with enzymes at temperature of 20C to 38C (68F to 100F)
for 28 to 72 hours. Under these conditions, about 90 percent
of the syrup is converted to ethyl alcohol.
Over time, synthetic methods for the production of ethyl
alcohol were developed. In one such method, ethylene
(ethene; CH2=CH2) is treated with sulfuric acid and water
to= obtain ethyl alcohol. That method was popular during
the 1950s and 1960s. Then, a new method for making the
compound was invented. In that process, ethylene and water
are heated together at high temperatures [300C to 400C
(570F to 750F)] and high pressures [1,000 pounds per
square inch (6.9 megaPascals)] over a catalyst of phosphoric
acid (H3PO4). The efficiency of this method is greater than

the older method, and there are fewer environmental consequences
from making ethyl alcohol by this process.
As of 2003, about 94 percent of all ethyl alcohol was
produced by fermentation. The remainder was produced by
the phosphoric acid method.

   In 2005, 10,500 million liters (2,790 million gallons) of
ethyl alcohol were produced by fermentation methods. Of
that amount, 92 percent was used as a fuel or an additive in
fuels. Many experts suggest that consumers use a mixture of
gasoline (90 percent) and ethyl alcohol (10 percent) called
gasohol as a vehicle fuel because it burns more completely
and releases fewer harmful byproducts to the environment.
Although gasohol has not yet become very popular in the
United States, it is widely used in some other parts of the
world, most notably, in Brazil.
Of the remaining 8 percent of ethyl alcohol produced by
fermentation, half was used in industrial operations, as a
solvent or intermediary in the preparation of other chemical

compounds; and half was used in the production of alcoholic
beverages.
In 2005, about 650 million liters (170 million gallons) of
ethyl alcohol were produced by the phosphoric acid method.
Of that amount, 60 percent was used for industrial solvents
in the manufacture of toiletries and cosmetics, coatings and
inks, detergents and household cleaners, pharmaceuticals,
and other products. The remaining 40 percent was used in
the preparation of other chemical compounds, including
ethyl acrylate, vinegar, ethylamines, ethyl acetate, glycol
ethers, and miscellaneous materials.
Ethyl alcohol commonly occurs in one of three general
forms. Absolute alcohol is ethyl alcohol that contains less than
1 percent impurities, such as water. Absolute alcohol is very
difficult to make because ethyl alcohol will absorb water from
the atmosphere or any other source that is available. The ethyl
alcohol used in fuels and almost all industrial operations is a
mixture of 95 percent ethyl alcohol and 5 percent water. Both
absolute and 95 percent ethyl alcohol are extremely toxic.
Ingestion of even very small amounts of either liquid has
serious health effects that may include death.
The alcohol with which most people commonly come into
contact is ethyl alcohol mixed with water in alcoholic beverages,
such as beer, wine, gin, vodka, rum, or bourbon. In such
beverages, the concentration of ethyl alcohol ranges from a
few percent to 50 percent.
The effects produced by ethyl alcohol on the human body
depend on the type of beverage consumed and the time
taken for consumption. Drinking a 5-percent beer over an
hour has a very different effect on the body than drinking a
50-percent vodka in five minutes.
Ethyl alcohol is a central nervous system depressant.
After ingestion, it passes through a person’s stomach and
the small intestine, where it is absorbed rapidly into the
bloodstream. It then travels throughout the body, interfering
with the normal functioning of the nervous system and
producing symptoms such as drowsiness, slurred speech,
blurred vision, unsteady gait, impaired judgment, and
reduced reaction time. With greater concentrations of alcohol
in the blood, these symptoms may become more severe,
resulting in coma and death.

Ascorbic Acid (The Vitamin C)

  Ascorbic acid (as-KOR-bik AS-id), or vitamin C, is one of
the most important dietary vitamins for humans because it
plays a crucial role in building collagen, the protein that
serves as a support structure for the body. It is a watersoluble
vitamin, which means that the body excretes any
excess vitamin C in the urine and cannot store a surplus.
For that reason, humans must consume vitamin C in their
daily diets. Vitamin C is found in many fruits and vegetables
and most kinds of fresh meat. Citrus fruits, such as oranges
and lemons, are especially rich in the compound.

Humans have known about the consequences of vitamin C
deficiency for centuries. People traveling long distances
on land or by sea often came down with an illness called
scurvy. The same illness struck people living in their own
homes during long winters. The disease was characterized by
pain and weakness in the joints, fatigue, bleeding gums,
tooth loss, slow healing of wounds, and bruising. These symptoms
were caused as the body’s connective tissue broke down

and small blood vessels ruptured. These symptoms began to
disappear as fresh foods became more available. If they did
not get enough fresh food in their diets, people could die of
scurvy.
Scurvy was common enough that many people searched
for its cause and cure. Sailors were especially vulnerable to

the disease, and the first recorded investigations involving
vitamin C were done by seafaring men. In 1536, French
explorer Jacques Cartier (1491–1557) cured his sailors of
scurvy by following the advice of Indians in Newfoundland,
feeding them extract of pine needles. Scottish physician
James Lind (1716–1794) began investigating the disease in
1747. He read many historical accounts of the diseases and
combined that information with his own observations to
deduce that scurvy occurred only among people with very
limited diets. He went on a ten-week sea voyage and fed the
solders various foods to see which ones were best at curing
scurvy. Citrus fruits proved to be most effective in preventing
the disease, a result that Lind reported in 1753. Captain
James Cook (1728–1779) led expeditions to the South Seas in
the late 1700s and kept his crew healthy by feeding them
sauerkraut. In 1795 the British navy began serving its sailors
a daily portion of lime juice, and two things happened: British
sailors stopped getting scurvy, and people began calling
sailors ‘‘limeys.’’
  Many people refused to believe that scurvy was caused by
a dietary deficiency, suggesting that it was instead the result
of eating bad food or lack of exercise. In 1907, Norwegian
biochemists Alex Holst (1861–1931) and Theodore Frohlich
conducted a study in which guinea pigs were fed an experimental
diet that caused them to develop scurvy. The link
between the vitamin and the disease was firmly established
by this research. Ascorbic acid was first isolated independently
by the Hungarian-American biochemist Albert Szent-
Gyo¨rgi (1893–1986) and the American biochemist Charles
Glen King (1896–1988) in 1932. It was synthesized a year
later by the English chemist Sir Walter Norman Haworth
(1883–1950) and the Polish-Swiss chemist Tadeusz Reichstein
(1897–1996), again working independently of each other.

  Plants and most animals (humans and guinea pigs being
two exceptions) synthesize vitamin C in their cells through a
series of reactions in which the sugar galactose is eventually
converted to ascorbic acid. For many years, the compound has
been made commercially by a process known as the Reichstein
process, named after its inventor Tadeusz Reichstein. This
process begins with ordinary glucose, which is converted to
another sugar, sorbitol, which is then fermented to obtain

yet another sugar, sorbose. The sorbose is then converted
step-by-step into a series of other products, the last of which
is ascorbic acid.
  Chemists have long been searching for an alternative to
the Reichstein process because it uses so much energy and
produces by-products that are hazardous to the environment.
In the 1960s, Chinese scientists developed a method
that involves only two steps in the synthesis of ascorbic
acid, and in the early 2000s, Scottish scientists were
attempting to develop a method that involved only a single
step using fermentation. Currently, however, the Reichstein
process remains the most popular method for making the
compound.

  The best known use of vitamin C is as a nutritional
supplement, taken to ensure that one receives his or her
daily minimum requirement of the vitamin. The recommended
daily allowance (RDA) of vitamin C for adults is 60
milligrams per day. Anyone who eats a well-balanced diet
that includes citrus fruits, tomatoes, and green leafy vegetables
probably does not need to take a vitamin supplement.
However, the amount of vitamin C one normally receives
from a supplement is unlikely to cause any harm.

  In addition to its nutritional uses, ascorbic acid has a
number of other industrial applications, including:
• As a food preservative;
• As a reducing agent in chemical processes;
• As a preservative in foods;
• As a color fixing agent in meats, helping meats keep
their bright red appearance;
• As an additive to bread dough, where it helps increase
the activity of yeast used in the dough; and
• As a treatment for abscission in citrus plants, the tendency
for a plant to lose its leaves, flowers, and fruits.

Petroleum

  Petrolatum (peh-tro-LAY-tum) is a mixture, not a compound.

Mixtures differ from compounds in a number of

important ways. The parts making up a mixture are not

chemically combined with each other, as they are in a compound.

Also, mixtures have no definite composition, but

consist of varying amounts of the substances from which

they are formed.

   Petrolatum is a complex mixture of hydrocarbons

derived from the distillation of petroleum. Hydrocarbons

are compounds that contain only carbon and hydrogen. The

hydrocarbons that make up petrolatum belong to the

methane (saturated or alkane) family of hydrocarbons with

the general formula CnH2N+2. Some members of the family

include methane (CH4), ethane (C2H5), propane (C3H8), and

butane (C4H10).

   Petrolatum occurs in a semi-solid or liquid form. The

semi-solid form is also called petroleum jelly or mineral jelly

and is commercially available under a number of trade

names, including Kremoline, Pureline, Sherolatum, and

VaselineTM. It ranges in color from white to yellowish to

amber. It is practically odorless and tasteless. It melts over a

wide range, from about 38 C to about 55 C (100 F to 131 F).

The liquid form is also known as liquid paraffin, mineral

oil, or white mineral oil. Such products are sold commercially

under trade names such as Alboline, Drakeol, Frigol,

Kremol, and Paroleine. It is a colorless, tasteless, and odorless

oily liquid.

   

   Oil was first discovered in the United States in the 1850s

in western Pennsylvania. A chemist from Brooklyn, New

York, Robert Augustus Chesebrough (1837–1938), visited

the new wells and noticed a wax-like material sticking

to the petroleum drilling rods. He learned that oil workers

used the ‘‘rod wax’’ to heal burns on their skin. Chesebrough

eventually extracted and purified the substance—petrolatum—

from petroleum and began manufacturing it in 1870. He

received several patents for his discovery and in 1878, he

gave his product the trade name of VaselineTM. His product

quickly became popular as an ointment for wounds and

burns. Unlike the animal and vegetable oils then being used

for that purpose, petrolatum did not spoil. By the late 1870s,

VaselineTM was selling at the rate of one jar everyminute in

the United States. In 1880, it was added to the U.S. Pharmacopoeia,

a manual that lists drugs used in medical practice.

   Petrolatum is a product of the fractional distillation of

crude oil. Crude oil is a complex mixture of hundreds or

thousands of compounds. These compounds can be separated,

or distilled, from each other by heating crude oil to high

temperatures. As the temperature of the crude oil rises,

various groups or a ‘‘fraction’’ of compounds boil off. The first

group of compounds includes gaseous compounds dissolved

in crude oil. The next group of compounds includes compounds

with low boiling points. The next group of compounds

includes compounds with slightly higher boiling

points. And so on. Eventually, a tar-like mass of compounds

with very high boiling points is left behind in the distilling

tower. This residue is heated to separate liquids from solids

remaining behind. Some of these liquids and solids make up

the semi-solid and liquid forms of petrolatum.

  

   Petrolatum has a wide variety of uses, ranging from

personal care and medical applications to industrial uses.

The solid form, such as VaselineTM is used as a topical ointment

for the treatment of dry, cracked skin and to reduce the

risk of infection. It works as a moisturizing agent because it

reduces water loss from the skin, It helps prevent infection

because it creates a barrier over wounds that prevents disease-

causing organisms from entering the body. Solid petrolatum

is also an ingredient in many skin care and cosmetic

products, such as skin lotions, body and facial cleansers, antiperspirants,

lipsticks, lip balms, sunscreens, and after-sun

lotions. In hair products, it helps smooth frizzy hair by

allowing hair to retain its natural moisture. The formation

used in most of these products remains virtually unchanged

from that developed by Robert Chesebrough in the 1800s.

  

  Solid petrolatum is also used in industrial applications

for a variety of purposes, such as:

• As a softener in the production of rubber products;

• In the food processing industry, to coat raw fruits and

vegetables and to help products retain moisture;

• As a defoaming agent in the production of beet sugar

and yeasts;

• For the lubrication of firearms and machine parts;

• In the production of modeling clays;

• In the manufacture of candles, to prevent a candle from

shrinking as it cools after being burned;

• In the preparation of shoe polishes; and

• As an ingredient in rust preventatives.

   The primary use of liquid petrolatum is as a laxative, a

product that loosens the bowels. It also has a number of other

applications, such as an additive in foods such as candies,

confectionary products, and baked goods; as an ingredient in

personal care products, such as baby oil creams, hair conditioning

lotions, and ointments; in many different kinds of

pharmaceutical preparations; in the production of industrial

lubricants; as a softening agent in the manufacture of rubber,

textiles, fibers, adhesives, and machine parts; as dust

suppressants; and as dehydrating agents for a number of

industrial processes.

Sodium Hydroxide

  Sodium hydroxide (SO-dee-um hye-DROK-side) is a white
deliquescent solid commercially available as sticks, pellets,
lumps, chips, or flakes. A deliquescent material is one that
absorbs moisture from the air. Sodium hydroxide also reacts
readily with carbon dioxide in the air to form sodium carbonate.
Sodium hydroxide is the most important commercial caustic. A
caustic material is a strongly basic or alkaline material that
irritates or corrodes living tissue. The compound ranked number
11 among chemicals produced in the United States in 2004.

  Sodium hydroxide is produced commercially simultaneously
with chlorine gas by the electrolysis of a sodium
chloride solution. In this process, an electric current breaks
down sodium chloride into its component elements, sodium
and chlorine. The chlorine escapes as a gas, while the sodium
metal form reacts with water to form sodium hydroxide

2NaCl ! 2Na + Cl2
2Na + 2H2O ! 2NaOH + H2

  Sodium hydroxide can also be produced easily by means
of other chemical reactions. For example, the reaction
between slaked lime (calcium hydroxide; Ca(OH)2) and soda
ash (sodium carbonate; Na2CO3) produces sodium hydroxide:
Ca(OH)2 + Na2CO3 ! 2NaOH + CaCO3
None of these alternative methods can compete economically,
however, with the preparation by electrolysis.

Sodium hydroxide has a great variety of household and
industrial uses. It is the active ingredient in drain cleaners
such as Drano because it breaks up and dissolves the greasy
mass that is responsible for drain blockages. It is also an
ingredient in many other household products, including oven
cleaners, metal polishes, and hair straighteners. Sodium
hydroxide is also used in the preparation of homemade and
processed foods. It is used in the preparation of soft drinks,
chocolate, ice creams, caramel coloring, and cocoa. Hominy, a
starchy food similar to grits, is made by soaking corn kernels
in a solution of sodium hydroxide in water. Bakers glaze
pretzels and German lye rolls with a weak lye solution before
baking them. The lye gives baked goods a crisp crust. Some
people use lye to cure olives.
  The largest single use for sodium hydroxide is in the
production of organic compounds from which polymers are
made, such as propylene oxide and the ethylene amines, and
of the polymers themselves, including the polycarbonates
and epoxy resins. About a third of all the sodium hydroxide
produced in the United States goes to this application.
Another important use of sodium hydroxide is in the pulp
and paper industry, where it is used to digest (break down)
the raw materials from which pulp and paper are made.
About 13 percent of all the sodium hydroxide made in the

  United States goes to this application. Sodium hydroxide is
also an important raw material in the manufacture of soap.
The method by which soap is made has not changed very
much for thousands of years. A fat or oil is added to a boiling
solution of sodium hydroxide in water. The fat or oil hydrolyzes
into its component parts, glycerol and fatty acids. The
sodium hydroxide then reacts with the fatty acids, forming
sodium salts. The sodium salt of a fatty acid is a soap. Sodium
hydroxide is also an important raw material in the manufacture
of inorganic compounds, especially sodium and calcium
hypochlorite, sodium cyanide, and a number of sulfur-containing
compounds. Some other important uses of sodium
hydroxide include:

• In the manufacture of cellophane and rayon;
• As a neutralizing agent during the refining of petroleum;
• In the manufacture of aluminum metal;
• For the refining of vegetable oils;
• As an agent for peeling fruits and vegetables for processing;
• In the extraction of metals from their ores;
• For the processing of textiles;
• In water treatment facilities;
• For etching and electroplating operations; and
• In a wide variety of research laboratory applications.

  Sodium hydroxide is one of the most caustic substances
known and a strong irritant to the skin, eyes, and respiratory
system. Exposure to sodium hydroxide dust, powder, or solid
can cause burning of the skin and eyes, with possible permanent
damage to one’s vision. Ingestion of the compound

causes burning of the mouth, esophagus, and stomach, resulting
in nausea, diarrhea, internal bleeding, scarring, and permanent
damage to the lungs and gastrointestinal system.
More serious results, such as a drop in blood pressure and
collapse, are also possible.

Caffeine

  Caffeine (kaf-EEN) is an organic base that occurs naturally
in a number of plant products, including coffee beans,
tea leaves, and kola nuts. It occurs as a fleecy white crystalline
material, often in the form of long, silky needles. It
usually exists as the monohydrate, C8H10N4O2 H2O, although
it gives up its water of hydration readily when exposed
to air.
  Scientists believe that humans have been drinking beverages
that contain caffeine for thousands of years. The first
recorded reference to a caffeine drink can be found in a
Chinese reference to the consumption of tea by the emperor
Shen Nung in about 2700 BCE. Coffee is apparently a much
more recent drink, with the earliest cultivation of the coffee
tree dated at about 575 CE in Africa.
  Caffeine was first studied scientifically by two French
chemists, Joseph Bienaime´ Caventou (1795–1877) and Pierre
Joseph Pelletier (1788–1842), who were very interested in
the chemical properties of the alkaloids. Between 1817
and 1821, Caventou and Pelletier successfully extracted
caffeine, quinine, strychnine, brucine, chinchonine, and
chlorophyll (not an alkaloid) from a variety of plants. The
first synthesis of caffeine was accomplished in 1895 by the
German chemist Emil Hermann Fischer (1852–1919), who was
awarded the 1902 Nobel Prize in chemistry for his work on
the alkaloids.

  Caffeine belongs to a class of alkaloids called the methylxanthines.
Chocolate, from the cocoa tree Theobroma cacao
contains another member of the class, theobromine. Both
caffeine and theobromine are stimulants, that is, compounds
that act on the nervous system to produce alertness, excitement,
and increased physical and mental activity.
  Caffeine can be extracted from coffee, tea, and kola
plants by one of three methods. These methods are used
primarily to produce the decaffeinated counterparts of the
products: decaffeinated coffee, decaffeinated tea, or decaffeinated
soft drinks. A commercial variation of these procedures
is to treat the waste products of tea or coffee processing, such
as the dust and sweepings collected from factories, for the
extraction of caffeine.
  In the first of the three extraction methods, the natural
product (coffee beans, tea leaves, or kola beans) are treated
with an organic solvent that dissolves the caffeine from the
plant material. The solvent is then evaporated leaving behind
the pure caffeine. A second method follows essentially the
same procedure, except that hot water is used as the solvent
for the caffeine. A more recent procedure involves the use of
supercritical carbon dioxide for the extraction process.
  Supercritical carbon dioxide is a form of the familiar gas
that exists at high temperature and high pressure. It behaves
as both a liquid and a gas. Not only is the supercritical carbon
dioxide procedure an efficient method of extracting caffeine,
but it has virtually none of the harmful environmental and
health problems associated with each of the other two methods
of extraction.
  Caffeine is also made synthetically by heating a combination
of the silver salt of theobromine (C7H8N4O2Ag) with
methyl iodide (CH2I), resulting in the addition of one carbon
and two hydrogens to the theobromine molecule and converting
it to caffeine.
  Caffeine is used in foods and drinks and for medical
purposes. Its primary action is to stimulate the central nervous
system. People drink coffee, tea, or cola drinks to stay
awake and alert because caffeine creates a feeling of added
energy. It does this by increasing heart rate, improving blood
flow to the muscles, opening airways to aid breathing, and
releasing stored energy from the liver to provided added fuel
for the body. In large quantities, caffeine can also cause
nervousness, insomnia, and heart problems. The effects of
caffeine can linger in the body for more than six hours. In
medical applications, caffeine is sometimes used as a heart
stimulant for patients in shock, to treat apnea (loss of breathing)
in newborn babies, to counteract depressed breathing
levels as a result of drug overdoses, and as a diuretic.
  Caffeine stimulates the brain in two ways. First, because
it has a chemical structure similar to that of adenosine, it
attaches to adenosine receptors in the brain. Adenosine is a
substance that normally attaches to those receptors, slowing
brain activity and causing drowsiness. By blocking those
receptors, caffeine increases electrical activity in the brain,
creating a feeling of alertness. Caffeine also works in the
brain like drugs such as heroin and cocaine, although in a
much milder way. Like those drugs, caffeine increases dopamine
levels. Dopamine is a chemical present in the brain that
increases the body’s feeling of pleasure.
Studies have shown that caffeine can become addictive.
People who use the compound eventually need to take more
and more of it to get the same effect. When some people try to
stop using caffeine, they may suffer from headache, fatigue,
and depression, though these symptoms can be controlled
by gradually reducing the amount of caffeine consumed.
Either way, withdrawal symptoms end after about a week.

Glucose

  Glucose (GLOO-kose) is a simple sugar used by plants and
animals to obtain the energy they need to stay alive and to
grow. It is classified chemically as a monosaccharide, a compound
whose molecules consist of five- or six-membered
carbon rings with a sweet flavor. Other common examples
of monosaccharides are fructose and galactose. Glucose
usually occurs as a colorless to white powder or crystalline
substance with a sweet flavor. It consists in two isomeric
forms known as the D configuration and the L configuration.
Dextrose is the common name given to the D conformation of
glucose.
Credit for the discovery of glucose is often given to the
German chemist Andreas Sigismund Marggraf (1709–1782).
In 1747, Marggraf isolated a sweet substance from raisins
that he referred to as einer Art Z cker (a kind of sugar) that
we now recognize as glucose. More than 60 years later, the
German chemist Gottlieb Sigismund Constantine Kirchhof
(1764–1833) showed that glucose could also be obtained from
the hydrolysis of starch and that starch itself was nothing
other than a very large molecule (polysaccharide) composed
of many repeating glucose units. The molecular structure
for glucose was finally determined in the 1880s by German
chemist Emil Fischer (1852–1919), part of the reason for
which he was awarded the 1902 Nobel Prize in chemistry.
  Glucose is synthesized naturally in plants and some single-
celled organisms through the process known as photosynthesis.
In this process, sunlight catalyzes the reaction
between carbon dioxide and water that results in the formation
of a simple carbohydrate (glucose) and oxygen. The overall
reaction can be summarized by a rather simple chemical
equation:
6CO2 + 6H2O ! C6H12O6 + 6O2
However, photosynthesis actually involves a number of
complex reactions that occur in two general phases, the light
reactions and the dark reactions.
Glucose is produced commercially through the steam
hydrolysis of cornstarch or waste products containing cellulose
(a large molecule composed of glucose units) using a
dilute acid catalyst. The product thus obtained is typically
not very pure, but is contaminated with maltose (a disaccharide
consisting of two molecules of glucose joined to each
other) and dextrins (larger molecules consisting of a number
of glucose units joined to each other).

  Glucose is the primary chemical from which plants and
animals derive energy. In cells, glucose is broken down in a
complex series of reactions to produce energy with carbon
dioxide and water as byproducts.
  Glucose also has a number of commercial uses, nearly all
of them related to the food processing business. It is used in
the production of confectionary products; chewing gum; soft
drinks; ice creams; jams, jellies, and fruit preparations; baby
foods; baked products; and beers and ciders. A relatively small
amount is used for non-food purposes, primarily in the production
of other organic chemicals, such as citric acid, the amino
acid lysine, insulin, and a variety of antibiotics.
The most important health problem associated with
glucose is diabetes. Diabetes is a medical condition that
develops when the body either does not produce adequate
amounts of insulin or cannot use that compound properly.
Insulin is a hormone that controls the metabolism of glucose
in the body. If glucose is not metabolized properly, a
person’s body acts as if it is ‘‘starving.’’ Symptoms of diabetes
include excessive hunger, weight loss, and exhaustion.
If left untreated, the condition can result in coma
and death. Diabetics must have an artificial source of insulin
(usually from injections) and watch their diets to keep
these symptoms under control.

Urea

Urea (yoo-REE-uh) is a white crystalline solid or powder
with almost no odor and a salty taste. It is a product of the
decomposition of proteins in the bodies of terrestrial animals.
Urea is produced in the liver and transferred to the
kidneys, from which it is excreted in urine. The compound
was first identified as a component of urine by French chemist
Hilaire Marin Rouelle (1718–1799) in 1773. It was first
synthesized accidentally in 1828 by German chemist Friedrich
Wo¨hler (1800–1882). The synthesis of urea was one of the
most important historical events in the history of chemistry.
It was the first time that a scientist had synthesized an
organic compound. Prior to Wo¨hler’s discovery, scientists
believed that organic compounds could be made only by the
intervention of some supernatural force. Wo¨hler’s discovery
showed that organic compounds were subject to the same set
of natural laws as were inorganic compounds (compounds
for non-living substances). For this reason, Wo¨hler is often
called the Father of Organic Chemistry.

The formation of urea is the evolutionary solution to the
problem of what to do with poisonous nitrogen compounds
that formed when proteins decompose in the body. Proteins
are large, complex compounds that contain relatively large
amounts of nitrogen. When they decompose, that nitrogen is
converted to ammonia (NH3), a substance that is toxic to
animals. If animals are to survive the decomposition of proteins
(as happens whenever foods are metabolized), some
method must be found to avoid the buildup of ammonia in
the body.

  That method involves a series of seven chemical reactions
called the urea cycle by which nitrogen from proteins

is converted into urea. Although high concentrations of urea
do pose a risk to animal bodies, the urea formed in these
reactions is normally excreted fast enough to avoid health
problems for an animal.
Urea is produced commercially by the direct synthesis
of liquid ammonia (NH3) and liquid carbon dioxide (CO2).
The product of this reaction is ammonium carbamate
(NH4CO2NH2):
2NH3 + CO2 ! NH4CO2NH2
Ammonia and carbon dioxide do not react with each
other under normal conditions of temperature and pressure.
If the pressure is raised to 100 to 200 atmospheres (1750 to
3000 pounds per square inch) and the temperature is raised
to about 200C (400C), however, the reaction proceeds efficiently
with the formation of ammonium carbamate. When
the pressure is then reduced to about 5 atmosphere (80
pounds per square inch), the ammonium carbamate decomposes
to form urea and water:
NH4CO2NH2 ! (NH2)2CO + H2O

  Urea is the sixteenth most important chemical in the
United States, based on the amount produced annually. In
2004, the chemical industry produced 5.755 million metric
tons (6.344 million short tons) of urea. Almost 90 percent of
that output was used in the manufacture of fertilizers. An
additional 5 percent went to the production of animal feeds.
In both fertilizers and animal feeds, urea and the compounds
from which it is made provide the nitrogen needed by growing
plants and animals for their good health and survival.
The other major use of urea is in the manufacture of various
types of plastics, especially urea-formaldehyde resins and
melamine.
Urea is also used:
• In the production of personal care products, such as
hair conditioners, body lotions, and dental products;
• In certain pharmaceutical and medical products, such as
creams to treat wounds and damaged skin;
• As a stabilizer in explosives, a compound that places
limits on the rate at which an explosion proceeds;
• In the manufacture of adhesives;
• For the flame-proofing of fabrics;
• For the separation of products produced during the
refining of petroleum;
• In the production of sulfamic acid (HOSO2NH2), an
important raw material in many chemical processes;
• As a coating for paper products; and
• In the production of deicing agents.

Penicillin

  The penicillins (pen-uh-SILL-ins) are a class of antibiotic
compounds derived from the molds Penicillium notatum and
Penicillium chrysogenum. The class contains a number of
compounds with the same basic bicyclic structure to which
are attached different side chains. That basic structure consists of two amino acids, cysteine and valine, joined to each
other to make a bicyclic (‘‘two-ring’’) compound. The different
forms of penicillin are distinguished from each other by
adding a single capital letter to their names. Thus: penicillin
F, penicillin G, penicillin K, penicillin N, penicillin O, penicillin
S, penicillin V, and penicillin X. A number of other
antibiotics, including ampicillin, amoxicillin, and methicillin,
have similar chemical structures.

  Penicillin was discovered accidentally in 1928 by the
Scottish bacteriologist Alexander Fleming (1881–1995). Fleming
noticed that a green mold, which he later identified as
Penicillium notatum, had started to grow on a petri dish that
he had coated with bacteria. As the bacteria grew towards the mold, they began to die. At first, Fleming saw some promise
in this observation. Perhaps the mold could be used to kill
the bacteria that cause human disease. His experiments
showed, however, that the mold’s potency declined after a
short period of time He was also unable to isolate the antibacterial
chemical produced by the mold. He decided that
further research on Penicillium was probably not worthwhile.

As a result, it was not until a decade later that Penicillium’s
promise was realized. In 1935, English pathologist
Howard Florey (1898–1968) and his biochemist colleague
Ernst Chain (1906–1970) came across Fleming’s description
of his experiment and decided to see if they could isolate the
chemical product produced by Penicillium with anti-bacterial
action. They were eventually successful, isolating and purifying
a compound with anti-bacterial action, and, in 1941, began trials
with human subjects to test its safety and efficacy (ability to kill
bacteria). The successful conclusion of those trials not only
provided one of the great breakthroughs in the human battle
against infectious diseases, but also won for Florey, Chain, and
Fleming the 1945 Nobel prize for Physiology or Medicine.

Penicillins are classified as biosynthetic or semisynthetic.
Biosynthetic penicillin is natural penicillin. It is produced by culturing molds in large vats and collecting and
purifying the penicillins they produce naturally. There are
six naturally occurring penicillins. The specific form of penicillin
produced in a culturing vat depends on the nutrients
provided to the molds. Of the six natural penicillins, only
penicillin G (benzylpenicillin) is still used to any extent.
  Semi-synthetic penicillins are produced by making chemical
alterations in the structure of a naturally occurring
penicillin. For example, penicillin V is made by replacing the
-CH2C6H5 group in natural penicillin G with a -CH2OC6H5
group.

  Penicillins are prescription medications used to treat a
variety of bacterial infections, including meningitis, syphilis,
sore throats, and ear aches. They do so by inactivating an
enzyme used in the formation of bacterial cell walls. With the
enzyme inactivated, bacteria can not make cell walls and die
off. Penicillins do not act on viruses in the same way they do
on bacteria, so they are not effective against viral diseases,
such as the flu or the common cold.
A number of side effects are related to the use of penicillin.
These side effects include diarrhea, upset stomach, and
vaginal yeast infections. In those individuals who are allergic to penicillins, side effects are far more serious and include
rash, hives, swelling of tissues, breathing problems, and
anaphylactic shock, a life-threatening condition that requires
immediate medical treatment.
Penicillin may alter the results of some medical tests,
such as those for the presence of sugar in the urine. Penicillin
can also interact with a number of other medications,
including blood thinners, thyroid drugs, blood pressure
drugs, birth control pills, and other antibiotics, in some cases
decreasing their effectiveness.
  Once promoted as wonder drugs, the use of penicillins
has declined slowly because of the spread of antibiotic resistance.
Antibiotic resistance occurs when new strains of bacteria
evolve that are resistant to existing types of penicillin.
One reason that antibiotic resistance has become a problem
is the extensive and often unnecessary use of penicillins.
When they are prescribed for colds and the flu, for example,
they have no effect on the viruses that cause those diseases,
but they encourage the growth of bacteria more able to
survive against penicillins.